Abstract

Have you ever wondered how nutritionists know how many Calories a certain food contains? In this project you will learn a method for measuring how many Calories (how much chemical energy) is available in different types of food. You will build your own calorimeter to capture the energy released by burning a small food item, like a nut or a piece of popcorn. This project gives a new meaning to the phrase "burning calories!"

Objective

The goal of this science project is to determine which food items store more chemical energy; you will determine this by burning the items and capturing the heat given off in a homemade calorimeter.

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Credits

Andrew Olson, Ph.D., and Sabine De Brabandere Ph.D., Science Buddies

Sources

USC Biology Department, 2004. Burning Calories: The Energy in Food.
Biology Department, University of Southern California. Out of print.

APA Style

Science Buddies Staff.
(2018, September 21).
Burning Calories: How Much Energy is Stored in Different Types of Food?
Retrieved from
https://www.sciencebuddies.org/science-fair-projects/project-ideas/FoodSci_p012/cooking-food-science/food-calorimeter

Last edit date: 2018-09-21

Introduction

You know that the energy that keeps your brain and body going comes from the food you eat. Your digestive system and the cells in your body break down the food and gradually oxidize the resulting molecules in a series of chemical reactions to release energy that your cells can use and store.

In this project you will learn a method for measuring how much chemical energy is stored in different types of food and express your result in Calories (note the capital "C"), as this is the unit of energy widely used to quantify food energy. To measure the chemical energy stored in food, you will oxidize the food much more rapidly than the cells in our body do by burning it in air. To do this, you will use a homemade bomb calorimeter that captures and measures the heat —the energy flow associated with differences in temperature— released by burning food. The basic idea of a calorimeter is to release all the stored energy at once and capture all the released heat with a reservoir of water. Measuring the temperature of the water at the beginning and at the end of the trial will allow you to calculate the energy put into heating up the water. As all this energy comes from the chemical reaction, the calculation will reveal the actual amount of energy released during the reaction, or the energy originally stored in the food. However, this is only true if all the energy released during the reaction is used to heat the water, and none of it gets "lost."

Bomb calorimeters that are used by scientists are made such that practically all of the energy released during the chemical reaction (like burning food) is captured by the water. Your homemade version will not reach the same efficiency to catch the heat released while burning the food; in other words, only a fraction of energy stored in the food will be converted to energy captured in the water and measured by your calorimeter. For example, some of the energy might go into heating up the surrounding air instead of the water. Even catching half of the energy released (an efficiency of 0.5) is acceptable for a homemade calorimeter, as it is very difficult to transfer all the chemical energy stored in the food and released during oxidation into the water in the calorimeter. Even with low efficiencies, this project will allow you to rank different kinds of food from more caloric to less caloric and will allow you to predict, with reasonable accuracy, the ratio of caloric content of different types of foods.

Now, let us see how to calculate the energy stored in the water for a measured increase in the temperature (in °C). The temperature difference times the mass of the water (in grams) will give you the amount of energy captured by the calorimeter, in calories, a unit of chemical energy. We can write this in the form of an equation:

[Please enable JavaScript to view equation]

where:

Qwater is the energy in the form of heat captured by the water, expressed in calories (cal);

mwater is the mass of the water, expressed in grams (g);

c is the specific heat capacity of water, which is 1 cal/(g °C) (1 calorie per gram per degree Celsius); and

(Tf - Ti) is the change in temperature, or the final temperature of the water minus the initial temperature of the water, expressed in degrees Celsius (°C).

The unit calorie (cal) (lowercase "c") is defined by the heat capacity of water. One calorie is the amount of energy that will raise the temperature of 1 g of water by 1 °C. When we talk about food energy, we also use the word Calorie (Cal) (note uppercase "C"), but it is a different unit. It is the amount of energy needed to raise the temperature of 1 kilogram (kg) (which equals 1,000 g) of water by 1 °C. So one Calorie (abbreviated as 1 Cal) is the same as 1,000 calories, also called 1 kilo calorie (kcal). In this project, for food Calories, we will be careful always to use an uppercase "C". Can you verify that the specific heat capacity of water equals 0.001 Cal/(g °C)?

Now we will work through an example to make sure that the equation is clear. (We will use made-up numbers for the example; you will have to do the project for yourself to get actual measurements.) So let us say that we start out with 100 milliliters (mL) of water in the calorimeter. Since 1 mL of water has a mass of exactly 1 g, this water will have a mass of 100 g (mwater = 100 g). The initial temperature of the water is 20 °C. After burning up a small piece of food, we measure the water temperature again, and find that the final temperature is 24 °C. Now we have all of the information we need to calculate the amount of heat captured by the calorimeter:

Can you see why the specific heat capacity of water has such strange units (cal/(g °C))? Notice that the grams (g) from the mass of the water and the degrees Celsius (°C) from the change in temperature cancel out with the grams (g) and degrees Celsius (°C) in the denominator of the units for specific heat. That way you are left with units of calories, a measure of energy, which is what you want.

Eating a balanced diet is fundamental to good health. This project will give you a chance to learn about how much energy your cells can extract from different types of food. It is important to remember, though, that energy is only one measure of nutritional value. As you are doing your background research on this project, try to find out about other measures of a balanced diet in addition to food energy.

Terms and Concepts

To do this project, you should do research that enables you to understand the following terms and concepts:

Calorie (Cal)

Oxidation

Calorimeter

Efficiency

Kilocalorie (kcal) and calorie (cal)

Questions

The reference level for a normal diet is 2,000 Calories. How many calories is this?

What are the basic chemical structures of fats, sugars (also called carbohydrates) and proteins?

How do these types of molecules differ in the amount of energy they contain?

Which of your food items do you think will release the most energy? Why?

What is meant by a "balanced" diet? Why is it important?

Bibliography

The U.S. Department of Agriculture is a good online source of information about nutrition. The links below are for general information, key nutritional recommendations, and special pages with information for kids:

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Digital scale with 0.1 g increments. A digital scale that would be suitable
(the Fast Weigh Digital Pocket Scale) is available from
Amazon.com.

Food items to test (dry items with a relatively high oil content and air trapped in them will generally work better), for example:

Cashew nuts, peanuts, or other whole nuts

Pieces of popcorn

Marshmallows

Croutons

Dry pet food

Churros

Croissant

Cheerios®

Lab notebook

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Recommended Project Supplies

Get the right supplies — selected and tested to work with this project.

Experimental Procedure

Safety Note: Adult supervision is required! As with any project involving open flame, there is a fire hazard with this project. Make sure you work on a non-flammable surface. Keep long hair tied back. Be careful handling the items used in this experiment as they may be hot! Wear safety glasses.

Assembling Your Calorimeter

Use the diagram shown in Figure 1, to guide you through building your calorimeter.

Figure 1. Diagram of homemade calorimeter.

Figure 1. Diagram of homemade calorimeter.

Use the can opener to remove the bottom from the larger can, so that you have a cylinder that is open on both ends. Be careful after opening as the edges might be sharp.

Note:If you are not using the kit, you will need to select two cans that nest inside one another. The smaller can needs to sit high enough so that you can place the cork, needle, and food item beneath it.

If your larger can is not aluminum on the inside, cover its inside with aluminum foil. Folding the edge of the aluminum foil over the edge of the can will keep it in place.

As a helper holds the can in place for you, use a hammer and nail to make holes around one edge of the cylinder. Space the holes about 4–5 cm apart. This edge will become the bottom edge when the can is standing vertically. The holes are shown as black spots in Figure 1 (the holes are there to allow air to come in and sustain the flame).

Put aside the large can and pick up the small can. As a helper holds the small can in place for you, use a hammer and nail to punch holes on opposite sides of the smaller can, about 1–4 cm from the top (open end). These holes will be used to attach the support. Depending on the relative height of the cans, you can choose to punch bigger holes and to put the support through them, or smaller holes and attach the support on top of the can with metal wire as shown in Figure 1 and 2.

Figure 2. The rod can be attached on top of the small can with metal wire as shown here, or threaded through the holes.

Put the supporting rod in place and attach it firmly. Note: It is not a good idea to use glue, as the glue might melt when the calorimeter gets hot.

Grasp the three needles and push the blunt ends into the cork, as shown in Figure 3. You will impale the food to be tested on the sharp ends of the needles. If needed, you can cut the cork lengthwise and let it rest on the cut side to reduce its height.

Figure 3. A cork with needles will hold the food items to be burned.

To construct the calorimeter:

Place the aluminum pie pan on a heat resistant surface.

Put the cork with needles sticking up in the middle of the pan.

Place the larger can over the cork on the aluminum pie pan.

Hang the smaller can inside the big cylinder.

The final result is shown in Figure 4.

Figure 4. Top-down view of the assembled calorimeter.

The smaller can will hold the water to be heated by burning the food samples. Use the graduated cylinder to measure how much water fills the can about half-full. Note this value (expressed in milliliters) in your lab notebook.

Taking Measurements

Copy the following table in your lab notebook. It will help you take notes as you perform your trials. Note that the mass of the water used in your calorimeter is not listed in the table. We advise you keep it the same for all your trials.

Food Item

Trial #

Food: Mi(g)

Food: Mf(g)

Water: Ti(°C)

Water: Tf(°C)

Qwater(Cal)

Qwater for 1 g food (Cal/g)

Average Qwater for 1 g food (Cal/g)

Food item 1

1

2

3

Food item 2

1

2

3

Food item 3

1

2

3

Table 1. Table in which to record measured and calculated values.

Decide on the food items you would like to study. For each item on your list, you will perform three measurements (trials). It is a good idea to repeat measurements to ensure consistent results. For each trial, you will impale a few pieces of each type of food; using a few items (and not just one) will allow you to burn a larger mass.

For each trial, you will:

Start with your calorimeter disassembled.

Weigh the food items to be burned and record the mass in the column "Food: Mi".

Impale the food items on the needles. Make sure all items touch, as shown in Figure 5. This will allow the flame to go from one item to the next.

Stir the water in the small can and measure the initial temperature (Ti). Record this temperature in the column "Water: Ti".

Note: After you have used your calorimeter, the water and can might still be cooling. Wait until the water reaches the same temperature as the environment or measure just before you put the can on top of your burning food. Remember to note the temperature in your table.

Have your calorimeter pieces close at hand, and ready for use.

Place the cork with food already impaled in the middle of the aluminum pie pan and light the food items with the long matches. Try the following if you have trouble igniting the food:

A big flame on your match will help ignite the food. A slight breeze helps create bigger flames. If you cannot find a fire-safe place with a breeze, consider using a fan (at safe distance).

Be patient, some food items like nuts might take a while to catch fire.

Definitely use long matches. This will make it safer and easier to keep the food item in the flame for a longer period of time.

When at least one food item catches fire, place the large can around the cork, then carefully place the smaller can in place above the flame.

Allow the food item to burn itself out. Use smoke coming out of the top as an indicator to evaluate if the burning is still in process. Try the following if you have trouble keeping the fire alive in the calorimeter:

Some food items might keep burning when they were put in the calorimeter just smoldering, but others will need a real flame to keep the burning active in the calorimeter. Experiment a bit with whether the food items burn better with a vibrant flame. Figure 6 might give you some ideas.

Enlarge the holes at the bottom of your cylinder to allow more air to pass through.

If the food in your calorimeter burns well, you might opt to tent the calorimeter with aluminum foil so less heat is lost to the environment, leaving a small opening at the top to allow air circulation.

Shortly after the food stops burning, carefully stir the water and measure the final temperature (Tf). Make sure the thermometer has reached a steady level before recording the value.

When the burnt food item has cooled, carefully remove it from the needles and weigh the remains. Record your value in the column "Food: Mf". Ideally, all the food should have burned up. If it is not, you will correct for this during your analysis by subtracting the final mass from the initial mass.

This will complete one trial for this food item.

Figure 5. Food items held in place by needles are placed so they touch each other.

Figure 6. Some food items will need to be in flames before being put under the calorimeter.

Repeat step 3 for two additional trials of this food item.

Repeat steps 2 and 3 for the other food items you have on your list.

Analyzing Your Data

To analyze your data, you will first calculate the energy captured by the water for each trial. As explained in the Introduction, the energy captured by the water (Qwater) can be calculated from the mass of the water in your calorimeter (mwater), the change in temperature of the water (Tf - Ti) and c, the specific heat capacity of water, which is 1 cal/(g °C) or 1/1,000 Cal/(g °C) using this equation:
[Please enable JavaScript to view equation]

Use the data in your data table to calculate the heat captured by the water for each trial and record your result in the column labeled Qwater. Following example where you burn 1.1 g of almonds and start out with 150 milliliters (mL) of water in the calorimeter might make the equation clear. Since 1 mL of water has a mass of exactly 1 g, this water has a mass of 150 g
(mwater = 150 g). If initially the temperature of the water is 20.0°C, and after burning the nuts in the calorimeter we measure a water temperature of 33.3°C, then the change in temperature of the water (Tf - Ti) equals 13.3°C, and the heat captured by the calorimeter Qwater is (150 g × 0.001 Cal/(g °C) × 13.3°C) or 2.0 Cal.

The energy you just calculated (Qwater) reflects energy released by the total amount of food burned, or (Mf - Mi) grams of food burned. Calculate how much Qwater would be if 1 g of food was burned by dividing Qwater by the amount of food burned (Mf - Mi). We call this the energy per unit weight, and it is expressed in Cal/g.
[Please enable JavaScript to view equation]

In the above example, QAlmond, 1 g of food equals 2.0 Cal/1.1 g, or 1.81 Cal/g. Do the calculation and write your number down in the column "Qwater for 1 g food".

Average the energy per unit weight released per individual food item over all three trials and write your value in the last column of your data table.

Since in a homemade calorimeter, only part of the energy contained in the food and released during burning transfers into energy stored in the water, the measured values only reflect a fraction of the chemical energy contained in the food. Some of the energy will get lost to, for example, heating up the surrounding air. Still, you can learn a lot from your obtained values. Below are some ideas to get you started:

Order your food items from more caloric to less caloric.

Create a graph, listing the food items on the x-axis and the caloric content of 1 g of food on the y-axis. Remember to label your axes, and add the units and a title to the graph.

Normalize the caloric content of food items to the caloric content of one food item. To do this, choose one food item from your list (e.g. Cheerios) and fill in the values, as directed by the equations in Table 2.

Food Item

Average Q for 1 g of Food (Cal/g)

Average Caloric Content Normalized to the Caloric Content of Cheerios

Cheerios

QCheerios

1

Food item 1

Qitem 1

Qitem 1 / QCheerios

Food item 2

Qitem 2

Qitem 2 / QCheerios

Table 2. Table containing the average caloric values as measured for all food items and the normalized caloric content relative to the caloric content of Cheerios.

These relative values show you how much more or less caloric food items are with respect to the item chosen (here Cheerios). For example, if the normalized caloric content for nuts would be 2, it would inform you that these nuts contain double as many calories per gram of food than Cheerios do.

Create a graph listing the food items on the x-axis and the normalized caloric content on the y-axis. Remember to label your axes, and add the units and a title to the graph.

Do you think the amount of Calories you measured is likely to be higher or lower than the true value for each food item? Why? If you can, look up the caloric content of your food items per gram of food. Does this confirm your hypothesis?

The Introduction states that, unfortunately, in homemade calorimeters, part of the energy stored in the food does not make it to the water. Can you suggest areas for improvement?

If you like this project, you might enjoy exploring these related careers:

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Variations

If you have the package information revealing the caloric content of some food items you measured, calculate the efficiency of your calorimeter. Table 3, below, can help you. Look back at the Introduction if you need to refresh your memory on the term efficiency.

Food Item

Average Q for 1 g food (Cal/g)

Caloric content for 1 g of food as shown on the package

Efficiency of the Calorimeter

Food item 1

Qitem 1

Qitem 1 -package

Qitem 1 / Qitem 1 -package

Food item 2

Qitem 2

Qitem 2 -package

Qitem 2 / Qitem 2 -package

Table 3. Table containing the average caloric values, as measured, and the package-listed caloric content, allowing you to calculate the efficiency of the calorimeter.

Is the efficiency for all food items more or less the same? If so, can you use the efficiency to calculate the real caloric content of a food item for which you do not have package information from your measurements?

Do background research to find out the approximate proportions of the different basic food chemicals (fats, carbohydrates, proteins) in each of the food items you tested. Can you draw any conclusions about the relative amounts of energy available in these different types of chemicals?

Do background research to find out the chemical composition of candle wax (paraffin). Design an experiment to determine the amount of energy released per gram of candle wax.

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